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The Development of the Periodic Table

The Development of the Periodic Table. Chapter 7 Section 1. Timeline of Development…. Ah Ha! My life has purpose again. 1790’s Antoine Lavoisier: compiled a list of elements (about 23) Mid-1800’s Scientists developed a way to determine atomic mass 1870 About 70 known elements.

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The Development of the Periodic Table

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  1. The Development of the Periodic Table Chapter 7 Section 1

  2. Timeline of Development… Ah Ha! My life has purpose again • 1790’s • Antoine Lavoisier: compiled a list of elements (about 23) • Mid-1800’s • Scientists developed a way to determine atomic mass • 1870 • About 70 known elements

  3. Organization • Meyer, Mendeleev & Moseley • Mendeleev gets most of the credit • Organized by atomic mass (just as Newlands) but changed columns • Organized into columns with similar properties • Left blank spaces for places where he thought elements should be, but weren’t discovered yet • Table 7.1?

  4. Mendeleev’s Predictions

  5. Why not atomic #? • It was found that some of Mendeleev’s elements were incorrectly placed • Why didn’t he use atomic number instead of atomic mass? • Answer: atomic #’s weren’t discovered until the early 1900’s

  6. Moseley’s Adaptation • After Henry Moseley discovered protons (and atomic number) he changed the organization and fixed Mendeleev’s problems • Periodic Law: • Periodic repetition of chemical and physical properties of the elements when arranged by increasing atomic number

  7. Parts of the Periodic Table Columns = Groups (or families) Rows = Periods

  8. Sections of the PT Transition Elements Inner Transition Elements

  9. Other periodic tables…

  10. Why? • Why do things behave the way they do? • The best predictor/explanation of why elements react are found in: • Their # of electrons • The way their electrons are organized • The size of the atoms • How much they want electrons or how much they want to get rid of electrons

  11. Valence Electrons • Electrons in the outermost energy level of an atom • Core Electrons: all electrons that are not in the valence shell Na 1s22s22p63s1

  12. Shielding constant (# of non-valence electrons) Effective Nuclear Charge Nuclear Charge (# of protons) Electron Shielding • Positives & Negatives are attracted to each other • Effective Nuclear Charge: describes the pull on the electrons by the nucleus Zeff = Z - S

  13. Atomic Size • 50ml + 50ml = ? • Atomic Size • Atoms of different elements have different sizes • What happens to Zeff as we go down a group? As we go across a period?

  14. Atomic Radius • What is it?

  15. Atomic Radius Trend Increases Increases

  16. Increases Increases Atomic Radius Trend • Why? • 1) As you go down a group, principle energy levels are added • (n=1, n=2, n=3) • This increases the radius

  17. Increases Increases Atomic Radius Trend • Why? • 2) As you go across a period: • No energy levels are added • Protons are added

  18. Ionic Radius • Ions: • An atom that has an overall positive or negative charge • Examples: • Cl-1(Chlorine with 17 protons and 18 electrons) • Mg2+ (Magnesium with 12 protons and 10 electrons) • What happens to size when atoms do this?

  19. Ionic Radius Trend • Positive Atoms • To become positive, atoms lose electrons • What happens to size if you lose electrons? • Hint: You now have more positives pulling in less negatives Positive Nucleus

  20. Positive Nucleus Ionic Radius Trend • Negative Atoms • To become negative, atoms gain electrons • What happens to size if you gain electrons? • Hint: You now have more negatives pulling out

  21. Chapter 7 Test • Monday – January 7th • Development of the Periodic Table • Periodic Trends (what & why) • Atomic radius • Ionic radius • Ionization energy • Electron Affinity • Isoelectronic • Ions • Groups of the Periodic

  22. Comparing Atomic Size • Remember isoelectronic • When atoms have the same electron configuration, which one is bigger? • Example: a) Na+ b) F-1 c) O-2 O-2 > F-1 > Na+ Na = +11 F = +9 O = +8 Radius decreases with increasing nuclear charge (# of protons)

  23. Sample 7.6

  24. Ionization Energy • The energy required to remove an electron from an atom • 1st IE: Energy to remove the first electron Na  Na+ + e- • 2nd IE: Energy to remove the 2nd Na+ Na2+ + e- • 3rd IE, 4th IE etc…

  25. Trend in 1st Ionization Energy Increases Hard to steal electrons Increases Easy to steal electrons

  26. IE Equations & Energies • We show the change through an equation: Na  Na+ + e- E=+495 Na+  Na+2 + e- E= +4562 • Why is the 2nd IE so much bigger?

  27. Spikes in IE

  28. Sample 7.7

  29. Electron Affinity • The measure of how much an atom wants to gain an electron • For most atoms, energy is released when this happens • Delta E = negative

  30. Affinity vs Ionization • Ionization energy • Cl  Cl+ + e-DE = 1251 kJ/mol • Electron Affinity • Cl + e- Cl-DE = -349 kJ/mol More negative = more energy given off = more favorable

  31. Electron Affinity Fluorine has the most electron affinity Increases Increases

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